Chemistry IGCSE (Paper 2)

solid

solid

• particles are close + touching each other

• regular + repeating pattern of particle arrangement

• vibrate about fixed positions but don’t move apart

• stronger force between particles than in a liquid

• not compressible

liquid

• particles are close + touching each other - more spaced than solid

• irregular particle arrangement

• particles move around and slide past each other

• forces between particles not as strong as a solid

• no fixed shape (takes shape of its container)

• not compressible

gas

• particles far apart

• irregular particle arrangement

• particles move freely + collide with each other

• very weak forces between particles

• no fixed shape or volume

• compressible

solid → liquid

= melting

solid heated → particles get energy → vibrate more violently → at certain temp. particles have enough energy to break free from position

liquid → gas/vapour

= evaporating

liquid heated → particles get more energy → move faster, weakening + breaking bonds holding liquid together → at certain temp, particles have enough energy to break bonds

liquid → solid

= freezing

requires significant temp. decrease + occurs at specific temp, different for each substance

gas → liquid

= condensing

gas cooled → particles lose energy → when particles collide, don’t have enough energy to bounce back → group together to form liquid

solid → gas

= sublimation

only happens to a few solids

Potassium manganate + water

• Put potassium manganate(VII) at bottom of water beaker

• Purple colour slowly spreads out to fill beaker

• Particles of potassium manganate(VII) diffuse out among particles of water

• Random motion of particles in liquid causes purple colour to eventually be evenly spread out in water

Dilution: If you added more water to final solution of potassium manganate, particles would spread further + solution would be less purple

Ammonia + hydrogen chloride

• Aqueous ammonia (NH₃) gives off ammonia gas

• Hydrochloric acid (HCl) gives off hydrogen chloride gas

• Set up experiment as in diagram → white ring of ammonium chloride forms in tube

• NH₃ gas diffuses from one end of tube + HCl gas diffuses from other → form ammonium chloride when they meet

• Ring doesn’t form exactly in middle - forms nearest end of hydrochloric acid

• Because ammonia particles are smaller + lighter so diffuse through air quicker

Bromine gas + air

• Bromine gas: brown, strong-smelling

• Fill half a gas jar full of bromine gas + other half full of air - separate gases with glass plate

• Remove glass plate → brown bromine gas diffuses slowly through air

• Random motion of particles means that bromine will eventually diffuse right through air

Solute

substance being dissolved

Solvent

liquid that solute dissolves in

Solution

mixture of solute + solvent that doesn’t separate out

Saturated solution

solution where max amount of solute has been dissolved - no more solute will dissolve in solution

Solubility

ability of substance to dissolve solvent

measured in grams of solute per 100g of solvent

solubility of most substances increases with temp

Solubility curve

Can be used to see the effect of temperature on solubility. Shows amount of a solute which dissolves to produce a saturated solution at any given temp.

Solubility curves can be used to predict how much solute will form when we cool a hot solution down.

Investigating how temp affects solubility of ammonium chloride

1. Make saturated solution by adding excess of ammonium chloride to 10cm³ of water in boiling tube

2. Stir solution + put boiling tube in water bath at 25ᵒC

3. After 5 mins, check that all excess solid has sunk to bottom of tube + use thermometer to check solution is at 25ᵒC

4. Weigh empty evaporating basin, pour some solution into basin

5. Re-weigh basin + contents, then gently heat using Bunsen burner to remove water - left with pure ammonium chloride

6. Re-weigh evaporating basin + contents

7. Repeat steps 1-6 twice but with water bath at diff temps

8. Use masses to work out solubility at each temp

element

consist of one type of atom only
e.g. oxygen, copper

mixture

• material composed of 2+ elements/compounds

• physically mixed together

• no chemical bond

• properties of mixture are mixture of properties of separate parts

• e.g. air (mixture of several gases), crude oil (mixture of hydrocarbons, mostly liquids)

compound

• made up of atoms of 2+ different elements joined by chemical bonds

• properties often totally different from properties of original elements

• e.g. carbon dioxide is compound formed from chemical reaction, one C atom reacts with two O atoms to form molecule of carbon dioxide

Pure substance

• Made of single element/compound

• Has specific melting + boiling point
  e.g. pure ice melts at 0ᵒC, pure water boils at 100ᵒC

• Mixture not pure - will melt/boil gradually over range of temperatures

filtration

used to separate insoluble solid from a liquid/solution

1. Put filter paper in funnel and pour in mixture

2. Liquid part runs through paper, leaving behind solid residue

crystallisation

used to separate soluble solid from solution

1. Pour solution into evaporating dish + gently heat solution
   Some water will evaporate, solution becomes more concentrated

2. Once some water has evaporated/when crystals start to form, remove dish from heat + leave solution to cool

3. Salt should start to form crystals as it becomes insoluble in cold, high conc. solution

4. Filter crystals out of solution + leave in warm place to dry

Paper chromatography

used to separate dyes

1. Draw line near bottom of filter paper (use pencil as pencil marks are insoluble so won’t dissolve in solvent)

2. Add spots of diff inks to the line at regular intervals

3. Loosely roll sheet up + put in beaker of solvent e.g. water

4. Ensure level of solvent is below baseline - don’t want inks to dissolve in solvent

5. Put lid on container to stop solvent evaporating

6. Solvent seeps up paper, carrying inks with it

7. Each dye in inks moves up paper at diff rate + forms spot in diff place

8. When solvent has nearly reached top of paper, take paper out of beaker + leave to dry

9. End result is called chromatogram

How chromatography separates mixtures

• Different dyes move up paper at different rates

• Some stick to paper, others dissolve more readily in solvent + travel quicker

• Distance travelled by dyes depends on solvent + paper used

Rf value

Rf = distance travelled by solute/distance travelled by solvent

1. To find distance travelled by solute, measure from baseline to centre of spot

2. Chromatography often used to see if certain substance is in mixture
   Run a pure sample of substance you think might be in mixture alongside sample of mixture itself
   If sample has same Rf values as one of the spots, they’re likely to be same

simple distillation

used to separate pure liquid from solution

1. Heat the solution
   Part of solution with lowest BP evaporates

2. Vapour is cooled, condenses + collected

3. Rest of solution is left behind in flask

4. Can use simple distillation to get pure water from seawater
   Water evaporates, condenses and is collected

Problem: can only be used to separate things with very different BPs

fractional distillation

used to separate mixture of liquids with different boiling points

1. Put mixture in flask + put fractionating column on top, then heat it

2. Different liquids have different BPs so evaporate at diff temps

3. Liquid with lowest BP evaporates first
   When temp on thermometer matches BP of liquid, it reaches top of column

4. Liquids with higher BPs also start to evaporate but column is cooler towards top, so they only get part of the way up before condensing + running back down towards flask

5. When first liquid has been collected, raise temp until next one reaches the top

atom

smallest particle of element
consists of electrons surrounding a nucleus that contains protons + neutrons

Sub-atomic particles

molecule

group of 2+ atoms chemically joined together

atomic number

number of protons in nucleus of atom

mass number

sum of number of protons + neutrons in nucleus of atom

isotopes

atoms of same element with same atomic number but different mass number

relative atomic mass

average mass of atom of an element
measured as ratio 1/12 of mass of atom of carbon-12

Calculating relative atomic mass

• Multiply % abundance of each isotope by its mass

• Add these numbers together

• Divide by total abundance (100%)

Periodic table

• Ordered in order of increasing atomic number

• Columns = groups

• Rows = periods

Elements in same group…

• have same number of electrons in outer shell

• have similar properties

• Properties of elements depend on number of electrons

• Number of electrons in outer shell is most important

Electron shell rules

Electrons occupy shells

1st shell: 2
2nd shell: 8
3rd shell: 8

Working out electronic configuration

• Period of element = number of shells containing electrons

• Group number = number of electrons in outer shell

• e.g. Sodium in period 3 so has 3 shells occupied
  First two shells must be full (2.8)
  In group 1 so has 1 electron in outer shell
  So electronic configuration = 2.8.1

Metals

• Elements on left of zigzag are metals

• Metals conduct electricity because they allow charge to pass through them easily

• Metal oxides are basic - they neutralise acids
  Metal oxides which dissolve form solutions with pH of 7+

Non-metals

• Elements on right of zigzag are non-metals

• Non-metals are poor electrical conductors

• Non-metal oxides are acidic - neutralise base
  Dissolve in water to form solutions with pH less than 7

Group 0

• Called noble gases, including helium, neon, argon

• Inert - don’t react much

• → takes a lot of energy to add/remove electrons from full outer shell of noble gas atom

Equations show…

Reactants + products of reaction
Can write word equations or chemical (symbol) equations

State symbols

(s) - solid

(l) - liquid

(g) - gas

(aq) - aqueous (dissolved in water)

Balancing chemical equations

• Put numbers in front of formulas where needed

• Find an element that doesn’t balance, write a number to try sort it out

• See where it gets you. May create another imbalance, if so, write another number and see where that gets you.

• Carry on correcting unbalanced elements until it solves

Relative formula mass

• Sum of relative atomic mass of all atoms

Mole

Unit for amount of substance

One mole of atoms/molecule of a substance has mass in grams equal to relative particle mass for that substance

Molar mass

Mass of one mole, measured in grams

Moles equation

Number of Moles = Mass in g / Mᵣ

Calculating masses in reactions

• Work out balanced equation

• Work out Mᵣ of reactant + product you’re interested in

• Find number of moles of the substance you know of

• Use balanced equation to find how many moles there’ll be of other substance

• Use number of moles to calculate mass

Percentage yield

• Calculate theoretical yield using balanced equation

• Percentage yield = actual yield/theoretical yield x 100

• 100% yield means you got all the product you expected

Empirical formula

Smallest whole number ratio of atoms in a compound

Calculating empirical formula

• List all elements in compound

• Write their experimental masses underneath

• Find number of moles of each element

• Turn numbers into ratio by dividing by smallest number of moles

• Get ratio in its simplest whole number form

Molecular formula

Actual number of atoms of each element in a compound

Calculating molecular formula

1. Find the mass of empirical formula

2. Divide molecular mass by formula mass

3. Multiply empirical formula by number obtained in step 2

Finding formulae using combustion

• Get crucible and heat until red hot (ensures that it’s clean and no traces of oil/water left)

• Leave crucible to cool, then weigh it, along with lid

• Add some clean magnesium ribbon to crucible
  Reweigh crucible, lid and magnesium ribbon

• Heat crucible containing magnesium
  Put lid on crucible to stop bits of solid escape, but leave small gap to allow oxygen to enter crucible

• Heat crucible strongly for around 10 mins

• Allow crucible to cool and reweigh crucible with lid + contents

• Use mass of magnesium oxide and initial mass of magnesium to calculate empirical formula

Finding empirical formulae using reduction

• Place rubber bung (with hole in middle) into test tube with small hole in end
  Weigh them using balance

• Take bung out of test tube + spread small amount of copper(II) oxide in middle of tube

• Re-insert bung + weigh test tube again

  Set up equipment as in diagram

• Expel air from test tube by gently turning on gas
  After 5 secs, light gas by holding burning splint next to hole in end of tube

• Use Bunsen burner to heat copper(II) oxide for 10 mins

• Turn off Bunsen burner + leave tube to cool

• Once tube has cooled, turn off gas and weigh tube with bung + contents

Salts

• All solid salts consist of lattice of +ve and -ve ions

• In some salts, water molecules are incorporated into lattice

• Water in lattice = water of crystallisation

• Solid salt containing water of crystallisation is hydrated

• If salt doesn’t contain water of crystallisation, it’s anhydrous

Calculating amount of water of crystallisation in a salt

• One mole of hydrated salt always has particular number of moles of water of crystallisation - formula shows how many

• e.g. hydrated copper sulfate has 5 moles of water for every one mole of salt
  So formula is CuSO₄.5H₂O (dot between CuSO₄ and 5H₂O)

• Many hydrated salts lose water of crystallisation when heated to become anhydrous

Concentration equation

Concentration (in mol/dm³) = Num of moles (mol) / Vol of solution (dm³)

Examples of finding conc. using experimental data

Molar volume

The space that one mole of a gas takes up

One mole of any gas always occupies 24dm³ (=24,000cm³) at room temp + pressure

Volume equation

Volume (dm³) = moles of gas x 24

Calculating volumes in reactions

• Calculate moles of substance you know of

• Find moles of gas using balanced equation

• Convert moles into volume using formula

Ions form…

when atoms lose/gain electrons
Negative ions (anions) form when atoms gain electrons
Positive ions (cations) form when atoms lose electrons

Using group number to predict ions formed

Groups 1, 2, 3 are metals. They lose electrons to form +ve ions.

Groups 5, 6, 7, are non-metals. They gain electrons to form -ve ions.

Elements in same group have same number of electrons in outer shell
So can lose/gain same number of outer electrons
So form ions with same charge

Silver

Ag⁺

Copper

Cu²⁺

Iron(II)

Fe²⁺

Iron(III)

Fe³⁺

Lead

Pb²⁺

Zinc

Zn²⁺

Hydrogen

H⁺

Hydroxide

OH⁻

Ammonium

NH₄⁺

Carbonate

CO₃²⁻

Nitrate

NO₃⁻

Sulfate

SO₄²⁻

Ionic compounds are produced by…

transfer of electrons

• When metal + non-metal react, metal atom loses electrons to form positive ion and non-metal gains these electrons to form negative ion

Formula of ionic compounds

Ionic compounds are made up of positively charged part + negatively charged part

Overall charge of ionic compound = 0

So negative charges must balance positive charges

Ionic dot and cross diagrams

• Dots represent electrons from one of the atoms

• Crosses represent atoms from the other atom

Ionic bond

Electrostatic attraction between oppositely charged ions

Giant ionic lattice

• Compounds with ionic bonding have giant ionic structures

• Ions held together in closely packed 3D lattice

• Electrostatic attraction between oppositely charged ions is very strong

• → a lot of energy needed to overcome strong attraction

• → high melting + boiling points

Ionic compound electrical conductivity

• Solid - don’t conduct electricity

• Molten/in aqueous solution - conduct electricity

Covalent bond

• Atoms make covalent bonds by sharing pairs of electrons with other atoms

• Each covalent bond provides 1 extra shared electron for each atom

Covalent bond is…

the strong electrostatic attraction between negatively charged pair of electrons and positively charged nuclei of the atoms involved

Dot and cross for diatomic molecules

Dot and cross for inorganic molecules

Dot and cross for organic molecules

Simple molecular substances

• Atoms within molecule are held together by very strong covalent bonds

• But forces of attractions between molecules are very weak

• → weak intermolecular forces = very low melting + boiling points, because molecules are easily separated

• Are gases/liquids at room temp or solid with low melting + boiling points

Molecules with high relative molecular mass…

have stronger intermolecular forces than smaller molecules

• Because there are more points along the larger molecules for intermolecular forces to act between them, so more energy needed to break forces

• → melting + boiling points of simple molecular substances increase as relative molecular mass increases

Giant covalent structures

• All atoms bonded to each other by strong covalent bonds

• Lots of bonds → takes lots of energy to break them

• → have very high melting + boiling points

• Don’t conduct electricity - even when molten (except for graphite)

Diamond

• Made of network of carbon atoms that each form four covalent bonds

• High melting point - strong covalent bonds take lots of energy to break

• Very hard - strong covalent bonds hold atoms in rigid lattice structure

• Doesn’t conduct electricity - no free electrons/ions

Graphite

• Each carbon atom forms three covalent bonds, creating layers of carbon atoms

• Soft + slippery - layers are held together weakly by intermolecular forces, so are free to slide over each other

• High melting point - covalent bonds in layers need lots of energy to break

• Conducts electricity - only 3 out of carbon’s 4 outer electrons are used in bonds, so each C atom has 1 delocalised (free) electron that can move

C₆₀ fullerene

• Hollow spheres made of 60 carbon atoms

• Not a giant covalent structure - made of large covalent molecules

• Soft - C₆₀ molecules only held by weak intermolecular forces so can slide over each other

• Poor electrical conductor - has 1 delocalised electron but electrons can’t move between molecules

Metallic lattice

Giant structure of positive ions surrounded by sea of delocalised electrons

Metallic bonding

Electrostatic attractions between nuclei of positive ions and sea of delocalised electrons

Metal electrical conductivity

Delocalised electrons can move through structure

→ metals can conduct electricity